1. Electrochemistry
Chapter 18

2. Redox Review: Oxidation numbers = oxidation states
Just like charges
Use periodic table as needed
Overall charge is 0 unless something indicated
Oxidation = loss of electrons; number gets more positive
Reduction = gain of electrons; number gets less positive
LEO GER
Half Reactions = separate the oxidation and reduction into equations
Balance separately and then combine

3. Balancing Redox BASE: For each half reaction:
Balance all elements not H and O
Balance O by adding H2O
Balancing H by adding H+
Balance charges by adding electrons
Need electrons to be equal so cancel out – multiply one half reaction or both by an integer to be equal
Add both half reactions together, cancel out anything same on both sides of arrow.
Add OH- ions to both sides of equation (to equal H+ ion)
combine these to make water
cancel out water from both sides if needed
ACID: For each half reaction:
Balance all elements not H and O
Balance O by adding H2O
Balancing H by adding H+
Balance charges by adding electrons
Need electrons to be equal so cancel out – multiply one half reaction or both by an integer to be equal
Add both half reactions together, cancel out anything same on both sides of arrow.

4. Examples

5. Electrochemistry
Study of the interchange of chemical and electrical energy
Generation of an electric current from a spontaneous reaction
Or a spontaneous reaction that generates an electric current
Galvanic cells
Device which changes chemical energy to electrical energy
Uses a spontaneous redox reaction to produce a current that can be used to do work
Electron transfer or flow of electrons causes current
Need connection for electrons to flow between half reactions
Separate the half reactions - one oxidation and one reduction

6. Components of Cell
Need two solutions which will house the half reactions
Oxidation = anode
Reduction = cathode
Electrons flow from anode to cathode
Solutions will be connected
Salt bridge – U tube filled with an electrolyte
Porous disk
Wires
Measure the flow of electrons – electrical potential or cell potential (Ecell)
Measured as voltage (V)
1 volt = 1 joule of work per columb ( 1V = 1 J/C)

7. Standard Reduction Potentials
p845 Table 18.1 – Standard Reduction Potentials
All based on hydrogen as standard potential = 0
All given as Reduction potentials (pay attention)
Each half reaction will be assigned a cell potential
To find the potential of the overall reaction, need to add both half reaction potentials together
E°cell = E°cathode – E°anode
When a half reaction is reversed, the sign of E° is reversed also
When the half reaction is multiplied by an integer, E° remains the same
Cell runs in the direction that gives a positive E value

8. Describing a Cell Line Notation – shorthand
Instead of drawing a cell diagram use this shorthand way
cathode on the right and the anode on the left.
phases of all reactive species are listed and their concentrations or pressures are given if not in standard states (1 atm. for gasses and 1M for solutions).
phase interfaces are noted with a single line ( | ) and multiple species in a single phase are separated by commas.
salt bridge or porous disk is shown as a double line ( || ).
EX: Mg (s) | Mg2+ (aq) || Al3+ (aq) | Al (s)
Pt (s) | H2 (g) | H+ (aq) || Cu2+ (aq) | Cu (s)

9. Describing a Cell Drawing a diagram: (Schematic diagram)
Label all parts
Show half reactions
Indicate direction of
electron flow

10. Describing a Cell
Write out or tell what is going on:

11. Relating free energy ∆G° = -nFE° Standard conditions
n = number of electrons
F = Faraday’s constant = 96,485 C/mol e-
∆G negative and Ecell positive = conditions for spontaniety

12. Batteries
Battery = galvanic cell, or group of galvanic cells connected
Potentials of individual cells added together to give battery potential
Types of batteries:
Lead storage battery – car battery
Contains lead (both cathode and anode) and sulfuric acid
Dry cell battery – AA, AAA, C, D
Acid: zinc (anode); carbon rod (cathode); moist paste MnO2 & NH4Cl
Alkaline: instead of NH4Cl has KOH or NaOH
Lithium-ion or Nickel-cadmium batteries
Fuel cells – reactants continuously supplied

13. Corrosion Involves oxidation of a metal prevent corrosion
Natural process for many metals
prevent corrosion
Galvanizing - coating metals to protect them from oxidation
Similar to plating
Alloying – mixture of metals
Cathodic protection – active metal attached by a wire to something, usually pipes

14. Electrolyisis Opposite of what did so far
Forcing an electric current through a cell to produce a chemical change
Cell potential would be negative
Non-spontaneous reaction
Requires an electrolytic cell (not galvanic cell)
Electrolysis of water
Electroplating